Chemistry 135—Unit 2: Molecular Bonding, Lewis Structures, and 3D Geometry
Lab preparation, logistics, and course logistics
Pre-lab questions for Lab 1: located in the lab manual; they tell you what to include in your lab notebook (e.g., a table of properties). There is an assignment in the sense that you must come to lab prepared; your TA will check your lab notebook.
Lab notebook expectations:
Write out the steps as given in the lab manual.
Use the word template recommended by the course (Word is best for formatting); UC students have access to the Microsoft suite.
Reports and questions for each lab are due before the start of the next lab period; submit online through Marcus.
The lab quiz will take place during lab.
Scheduling and alternating labs:
Bio and Chem labs can be scheduled in alternating weeks; e.g., Bio started Week 1; this week is Week Ending in 1; if you want to alternate on the same day, the sections should end in the same number.
There is an ongoing discussion about whether labs and content should be on the same or different days; the current understanding is that overlap should be avoided and scheduling is flexible.
Accessing links and resources:
If you can’t find the Zoom link, check the class resources tab or use the link for the 8 PM office hours (the same link is typically used).
The lab manual can be picked up in Lash Miller 204 (Chem club area), where Victoria College tutors also operate; corpus (course resources) will have details.
Quiz and content timing:
The quiz for Lab is not available before entering the lab; it is administered during lab.
Office hours and follow-up:
Office hours run after class, with time set aside for questions; questions asked in chat can be addressed there if not answered live.
What to expect today (Unit 2 introduction):
We will cover the first three topics today and save the next two for Thursday; the final lecture this week will begin discussion on GATS and GATS law (as noted in the session).
For readings, you should be at Chapter 3 at this point.
From quantum unit to molecule-building: overview and key concepts discussed
Transition in topics:
Move from quantum nature of the atom to building molecules (molecular orbitals, bonding).
We will learn how to apply those concepts to predict molecules and their properties.
Two important running themes:
How electrons populate atoms and molecules while obeying quantum rules.
How bond formation lowers energy and stabilizes structures via orbital overlap.
Visual cues used in session:
Energy diagrams to illustrate bond formation and bond energy.
Emphasis on the difference between atomic orbitals and molecular orbitals (MOs).
Ion formation and electron shielding concepts
Calcium example and ion formation:
Calcium tends to form a 2+ cation: Ca → Ca^{2+} + 2e^{-}.
Electron configuration: neutral Ca is
ext{Ca: } [Ar] \, 4s^{2}When it loses two 4s electrons, it becomes
ext{Ca}^{2+}: [Ar]Rationale: Atoms aim to achieve the closed-shell, noble gas electron configuration (like Argon in this case).
Shielding concept:
Inner core electrons shield valence electrons from nuclear charge to some extent.
Valence-shell electrons experience less repulsion from inner electrons; inner electrons effectively screen the nucleus.
Relevance when filling valence shells: adding electrons to the valence shell does not dramatically increase intra-shell repulsion because shielding mitigates it.
Practical takeaway:
When filling the valence shell by adding electrons, don’t overburden the shell with repulsion from inner electron screening; core electrons do shielding that affects effective nuclear charge felt by valence electrons.
Energy landscapes: H2 bond formation and basic energetic concepts
Energy vs internuclear distance (H–H example):
As two hydrogen atoms approach, their atomic orbitals overlap, electrons pair, and the system lowers its energy due to bond formation.
There is an energy minimum at the optimal bond distance (the bond length).
The energy difference between the minimum (bonded state) and the separated atoms is the bond energy.
If you pull the atoms apart beyond the bond length, you must input energy to break the bond; this is the bond energy magnitude.
Conceptual takeaway:
Bond formation corresponds to a favorable decrease in energy due to orbital overlap and electron sharing to satisfy valence needs.
The minimum on the energy curve defines the most stable bond length and the associated bond energy.
Molecular orbitals vs atomic orbitals; spin pairing
Molecular orbitals (MOs) arise from linear combinations of atomic orbitals (LCAO).
Atomic orbitals describe electrons in isolated atoms; MOs describe electrons in molecules where wavefunctions interfere and delocalize across the molecule.
Spin considerations in MOs:
When two electrons occupy the same molecular orbital, they must have opposite spins (Pauli exclusion principle).
This is why bonding MOs are typically filled with paired electrons in simple cases like H2.
Practical note:
The concept of molecular orbitals helps explain bonding strength and the distribution of electron density between atoms.
Electronegativity, bond polarity, and the continuum from covalent to ionic
Electronegativity and bond polarity:
Electronegativity determines how strongly an atom attracts bonding electrons.
In polar covalent bonds, bonding electrons are shared unequally, producing partial charges: δ− on the more electronegative atom and δ+ on the other.
Notation examples: δ− and δ+ (lowercase delta to indicate partial charges).
Bond polarity can be represented with arrows pointing from the less electronegative atom toward the more electronegative atom (electron density shifts in that direction).
Thresholds and practical rules of thumb:
A common rule-of-thumb threshold used in teaching is that a substantial ionic character tends to arise when the electronegativity difference Δχ is large (an approximate cut-off around Δχ ≈ 1.9 is often cited in this course context).
Bonds with Δχ around 1.9 can be considered highly polar covalent or approaching ionic in character, depending on context and other factors.
Examples discussed:
H–Cl bond: chlorine is highly electronegative, leading to a polar covalent bond with partial charges (δ− on Cl; δ+ on H). The bond is not ionic, but highly polar.
H–F and other halogen–hydrogen bonds were discussed in terms of polarity and electron density shifts.
Lithium fluoride (LiF) discussed as a classic ionic model: metals and nonmetals form ions (Li^+ and F^−) and attract electrostatically; the formation of ions does not necessarily require complete electron transfer in every case but tends to yield ionic solids when charges are localized.
Visual cues for polarity:
Polar bonds represented with arrow notation (toward more electronegative atom).
Partial charges shown with δ− and δ+ to indicate unequal electron sharing in covalent bonds.
Important caveat:
These are rules of thumb and there are exceptions; bond type exists on a continuum between purely covalent and ionic, influenced by context and molecular environment.
Examples and molecular polarity in practice
Typical polar covalent example discussed:
Hydrogen–chlorine (HCl) bond is polar covalent with significant partial charges, not fully ionic.
Ionic example discussed:
Lithium fluoride (LiF) is largely ionic due to large electronegativity difference and metallic vs nonmetal character.
Special case highlighted:
Some molecules can show very large electronegativity differences yet still exhibit covalent bonding in the solid due to lattice effects or bonding context; the teacher emphasizes that the 1.9 threshold is a guide, not a hard rule.
ClF3 example (guest lecturer note):
Chlorine trifluoride (ClF3) is used as a demonstration example for three-dimensional shape discussion and lone-pair placement; the molecule has five electron domains around Cl (three bonding pairs, two lone pairs) leading to a seesaw geometry in the classical VSEPR sense (not named in detail here beyond orientation discussion).
The facilitator notes not to memorize every shape name for regulatory purposes in certain spectroscopy contexts, but to understand the spatial arrangement for predicting intermolecular forces.
Lewis structures: a practical workflow with CO2 as an example
Purpose of Lewis structures: represent covalent bonds and lone pairs with a simple schematic.
Example: carbon dioxide, CO2
Step 1: Count total valence electrons
Carbon has 4 valence electrons; each oxygen has 6; total = $4 + 2\times6 = 16$ electrons.
Step 2: Choose central atom and draw initial bonds
Central atom: carbon (least electronegative among the considered atoms in CO2).
Draw a single bond from carbon to each oxygen (two bonds total), which uses 4 electrons (two bonds × 2 electrons per bond).
Step 3: Distribute remaining electrons as lone pairs to satisfy octets
After placing 4 electrons in bonds, 12 electrons remain to be placed as lone pairs.
Give full octets to the oxygens first: each oxygen gains 3 lone pairs (6 electrons per oxygen), using all 12 electrons.
Step 4: Check octets and adjust with multiple bonds if needed
Carbon currently has only 4 electrons around it (not an octet). To satisfy octets, move a lone pair from each oxygen to form a double bond with carbon, creating two C=O double bonds.
Result: CO2 with two C=O double bonds; each atom achieves a stable octet.
General takeaway:
This is a standard procedure for drawing Lewis structures: valence electrons counting, central atom selection, octet completion with lone pairs, and formation of multiple bonds as needed.
Valence Shell Electron Pair Repulsion (VSEPR) and three-dimensional geometry
Core idea: electron groups around a central atom arrange themselves to minimize repulsion to give predictable shapes.
What counts as an electron group:
Any bond (single, double, or triple) counts as one electron group.
Each lone pair counts as one electron group.
Shapes corresponding to electron groups around the central atom:
2 electron groups: linear geometry (180°)
3 electron groups: trigonal planar geometry (approx. 120°)
4 electron groups: tetrahedral geometry (109.5°)
5 electron groups: trigonal bipyramidal geometry (see-saw, T-shaped, or linear perspectives depending on lone pairs)
6 electron groups: octahedral geometry (90° and 180° interactions in certain axes)
How lone pairs influence observed molecular shapes:
Lone pairs occupy positions to maximize separation from bonding pairs, often favoring equatorial positions in trigonal bipyramidal geometries when lone pairs are involved.
For example, SF4 has a seesaw shape because it has five electron groups with one lone pair occupying an equatorial position, producing a distorted geometry for the four bonded atoms.
In XeF4 (not explicitly named in the transcript, but discussed as a relevant example later), two lone pairs in an octahedral electron geometry lead to a square planar molecular geometry.
Practical takeaway:
The angles listed (e.g., 180°, ~120°, ~109.5°) are idealized values assuming perfect geometries; real molecules can deviate slightly due to lone-pair repulsion and other effects.
Special guest annotation: xenon and chlorine trifluoride (ClF3) example
Xenon example (as discussed by the guest):
Xenon example with six electron groups and two lone pairs yields a square planar geometry for the molecule, illustrating how lone-pair placement influences the overall shape.
Chlorine trifluoride (ClF3) discussion:
The lecturer confirms chlorine trifluoride as the intended example (not carbon iodide).
ClF3 has five electron groups around chlorine: three bonding pairs and two lone pairs, leading to a seesaw molecular geometry.
Takeaway about shapes:
While it can be helpful to memorize shape names for certain topics, the emphasis here is on understanding the three-dimensional arrangement of electron groups and how lone pairs influence geometry, which matters for predicting intermolecular forces and reactivity.
Key takeaways on bonds, polarity, and practical modeling
Bonds exist on a continuum between covalent and ionic, influenced by electronegativity differences and context; the rules of thumb (e.g., Δχ ≈ 1.9 as a threshold) help guide expectations but are not absolutes.
Polar covalent bonds produce partial charges (δ−, δ+) and dipole arrows; ionic bonds involve significant electron transfer and electrostatic attraction between ions.
Lewis structures provide a practical way to count valence electrons, assign bonding and lone pairs, and assess octet satisfaction; they are stepping stones to predicting actual 3D shapes via VSEPR.
Molecular geometry and polarity interact to influence intermolecular forces (including hydrogen bonding in appropriate contexts) and properties like boiling point, solubility, and reactivity.
The course emphasizes the relationship between atomic-level interactions (orbitals, electron distribution) and observable properties of molecules, while acknowledging common exceptions and the need for experimental context.
Final reminders and announcements from the session
There is an office hours window after class (about one hour) for questions; use this time if you need more help.
The session encourages using the official word template for lab reports to minimize formatting issues; Word access is available to UC students.
For questions during class, use chat, and the instructor may address some questions live; if not, they will be addressed during office hours.
A guest lecturer (Anya Harlow) highlighted a xenon example and discussed lone-pair placement; also provided clarification about the ClF3 example and confirmed the naming to be chlorine trifluoride rather than chlorine triiodide.
The end of the session included a prompt to try a particular problem about ClF3; you are encouraged to attempt it and ask questions in office hours.
Quick reference formulas and conventions used in the notes
Electron configuration examples:
Neutral Ca: ext{Ca}: [Ar] \, 4s^{2}
Ca^{2+}: ext{Ca}^{2+}: [Ar]
Bond energy concept:
Bond formation lowers energy; the energy required to break a bond is the bond energy, often denoted as E_{ ext{bond}} with a positive value for bond breaking and a negative value for bond formation when considering the system’s energy change.
Valence electrons in Lewis structures:
Total valence electrons = sum of valence electrons from all atoms in the molecule.
VSEPR electron-group count examples:
2 groups → linear (180°)
3 groups → trigonal planar (≈120°)
4 groups → tetrahedral (≈109.5°)
5 groups → trigonal bipyramidal (see-saw, T-shaped, etc., depending on lone pairs)
6 groups → octahedral (various 90°/180° relationships depending on lone pairs)
Polarity indicators:
Dipole arrows point toward the more electronegative atom.
Denote partial charges with δ− and δ+ on the respective atoms.
Key precaution:
Treat thresholds like Δχ ≈ 1.9 as heuristic guides, not absolutes; real systems may vary due to competing effects and lattice environments.